Chemistry 2nd Semester Exam Review Sheet:
Chapters 9, 10, 11, 12, 13, 14, 15, and 16
Chapter 9: Chemical Quantities
1. Solve stoichiometry problems using grams molesmoles grams calculations. Examples below:
a. Calculate moles/grams of (NH4)2SO4 produced when 30.0 moles of NH3 react with H2SO4.
2NH3 + H2SO4  (NH4)2SO4
30.0 mol
x moles
30.0 moles x 1mol/2mol = 15.0 moles
15.0 moles (NH4)2SO4 x 132.2 grams/mole = 1983 = 1980 grams (to 3 s.f.)
b. Na2CO3 + 2HNO3  2NaNO3 + CO2 + H2O. How many moles/grams of sodium nitrate are produced by the
complete reaction of 100. grams of sodium carbonate? Moles of nitric acid needed to produce 500.0 grams of sodium
nitrate? Be able to calculate grams or moles of any reactant or product given a quantity of reactant or product.
Grams  Moles  Moles  Grams
100. g Na2CO3 x 1 mol/106.0 g x 2 mol NaNO3 / 1 mol Na2CO3 x 85.00 g/mol = 160.4 g
500.0 g NaNO3 x 1 mol/85.00 g x 2 mol / 2 mol x 63.02 g/mol = 370.7 grams
2. Calculate percent yield of products. Example below using (b) above.
What is the theoretical yield of NaNO3 if 100.0 grams of Na2CO3 are used in the reaction? What is the percent yield
of NaNO3 if 130.5 grams of NaNO3 are actually collected?
Theoretical yield = 160.4 grams (answer from above)
Percent Yield = (Actual Yield/Theoretical Yield) x 100 = (130.5 g/160.4 g) x 100 = 81.36%
3. What is a limiting reactant in a chemical reaction? The reactant that is used up first, thereby limiting the amount of
product that can be formed.
Chapter 10: Modern Atomic Theory
4. Describe Rutherford’s model of the atom. What could it not explain?
Nucleus (very small) consisting of protons with electrons in orbits around the nucleus. Could not explain how the
electrons remained in orbit.
5. Define electromagnetic radiation, wavelength, and frequency.
6. How do atoms gain energy to become excited? Absorb a photon of energy.
.How do they lose energy to return to their ground state? Release a photon of same energy as that absorbed.
7. What is the evidence that energy absorbed by the hydrogen atom is quantized? Specific wavelengths of light seen
in emission spectra for atoms (page 285-86)
8. Describe the Bohr model of the atom. How did it improve on Rutherford’s model?
Same as Rutherford with respect to nucleus. Electrons located on orbits at specific distances from nucleus.
9. Draw the orbital diagram (with the arrows) for an element: for example, phosphorus, sodium, and fluorine.
See book, notes, probs 54, 96
10. According to the wave-mechanical model, what are orbitals? Describe the shape of s and p orbitals.
Orbitals are regions of space around a nucleus where electrons have a probability of being found. S orbitals are
spherical, p orbitals are lobed. See sec 10.7 in text.
11. Identify elements from electron configurations. Ex: which elements are depicted below?
a. 1s22s22p1
b. 1s22s2
c. 1s22s22p63s23p2
boron
beryllium
silicom
12. From an electron configuration determine if an atom is in a ground or excited state. Ex: Which is an excited state?
a. 1s22s22p6
b. 1s22s22p53s2
c. 1s22s22p63s23p64s23d3
ground
excited
ground
(observe that the 2p electron has been promoted to the s sublevel)
13. Write the electron configurations for atoms and ions. Ex: a. K+
b. O2-
a. 1s2 2s2 2p6 3s2 3p6
d. 1s2 2s2 2p6 3s2
b. 1s2 2s2 2p6
c. 1s2 2s2
c.
Be
d. Mg
e. Ca
e. 1s2 2s2 2p6 3s2 3p6 4s2
Chapter 11: Bonding
14. Differentiate between ionic and covalent bonds. Ionic: electrons are transferred between a metal and nonmetal
forming ions (cation, anion). Covalent: electrons are shared between two nonmetals.
Which types of compounds are ionic and which are covalent. (see above)
15. Define electronegativity. Using the table on p. 320 in your textbook, arrange this set in order of increasing
electronegativity: Cl, Na, Al. (See pg 320) Na<Al<Cl
16. Which pairs are ionic, polar covalent, or covalently bonded? Mg-Cl, S-Br, O-O, C-O
See page 320-321
17. Compare/contrast the different geometric shapes of molecules. Determine the shape of different compounds from
their Lewis structures. Ex: Draw Lewis structures for H2S, NH3, SiCl4 and determine their molecular geometry.
Remember: calculate total number of valence electrons before drawing Lewis structures. Follow steps on page 342 and
rules on page 344. H2S is bent, NH3 is trig. pyramidal, SiCl4 is tetrahedral
Chapter 12: Gases
18. Write the names and formulas for all the gas laws. Derive all gas laws from PV = nRT, or use
P1V1/n1T1 = P2V2/n2T2 and discard the parts not in the problem.
19. Solve gas law problems. P. 390-393, do q. 7 a, 17 a, 33 a, 53 a, 85.
7a. 109.2kPa X 1 atm/101.325kPa = 1.078 atm
17a.
V1P1 = V2P2
P2 = V1P1/V2 = (117 mL)(652 mm Hg)/125mL = 610 mm Hg
33a.
V1/T1 = V2/T2
V2 = V1T2/T1 = (25mL) (273K)/298K = 23 mL
53a.
PV = nRT
V = nRT/P = (0.103 mol)(0.08206)(300K)/1.03 atm = 2.46 L
85. CaCO3  CaO + CO2
15.2 g CaCO2 x 1 mol/100.1 g x 1 mol/1 mol x 22.4 L/mol = 3.41 L
20. Explain the behavior of gases using the Kinetic Molecular Theory.
(see Sec 12.8, 12.9)
Chapter 13: Liquids and Solids
21. Compare/contrast solids, liquids, and gases. (with respect to distance between particles, movement of particles)
22. Draw and label the heating/cooling curve of water. (refer to your notes)
P. 418 do q. 18.
3.95 kJ/g x 26.98 g = 107 kJ/mole (molar heat of fusion for aluminum)
10.0 g x 1 mol/26.98 g = 0.371 moles of Al
0.371 moles x 107 kJ/mole = 39.7 kJ
23. Compare intramolecular with intermolecular forces. Define London dispersal forces, hydrogen bonds, and dipoledipole attractions. (see sec 13.3 in text)
Chapter 14: Solutions
24. What is a solution? Give examples. (see text, notes)
25. What is meant by solubility? Describe unsaturated, saturated, and supersaturated solutions.
(see notes)
26. Solve molarity, mass percent, and dilution problems. P. 447-449, do q. 17a, 33a, 45a, 55a, 70.
17a.
mass percent = mass solute/mass solution
Mass solute = mass solution x mass percent = 11.5 g x 0.0625 = 0.719 grams
33a.
M (molarity) = mol/L = 0.426 moles/0.500L = 0.852M
45a.
moles = M x L = 2.50 L x 13.1 M = 32.75 = 32.8 moles
32.8 moles x 36.46 g/mole = 1.20 x 103 g HCl
55a.
M1V1 = M2V2
M2 = M1V1/V2 = (0.105M)(0.425 L)/1.00 L = 0.0446 M
70,
do not need to do – we didn’t cover this material – not on final
Chapter 15: Acids and Bases
27. Define acids and bases according to Arrhenius and the Bronsted-Lowry model. List their basic properties.
28. Identify an acid and its conjugate base; a base and its conjugate acid. What is the relationship between acid
strength and conjugate base strength?
29. Compare a strong acid with a weak acid. What determines if an acid is a strong or a weak acid?
See your recent notes for this material. It’s all there.
30. Use the ion product constant of water ( Kw) to solve for [H+] and [OH-].
Kw = [H+] * [OH-]
Kw = 1.0 x 10-14
31. Find the pH and/or pOH of a solution with a given [H+] or [OH-] concentration, and the concentration of [H+] or [OH-]
from the pH or pOH.
pH = -log[H+]
[H+] = inverselog (-pH)
[OH-] = inverselog (-OH)
pH + pOH = 14
32. What happens when an acid reacts with a base? What happens when an acid reacts with metals?
Forms water and a salt. A single displacement reaction -- the metal cation replaces the H+ in the acid; H2 produced
33. A solution has a pOH of 12. Find [H+], [OH-], pH, and label if it’s acidic or basic.
pH = 14 – pOH = 2
[H+] = inverselog (-2) = 0.01 M
[OH-] = Kw/[H+] = 1.0 x 10-14 /0.01 = 1 x 10-12 M
Chapter 16: Equilibrium
34. Describe chemical equilibrium. (notes)
35. Describe conditions that affect reaction rates. (discuss concentration, temperature, and volume/pressure changes)
36. What is Le Chȃtelier’s Principle? How does it predict equilibrium shifts when a system at equilibrium is disturbed?
See all recent notes and WS.